Atomic Structure

Question 1

What is most of the volume of an atom made up of?

Answer 1

Electrons (nucleus is very small compared to the total size of the atom)

Question 2

Where is most of the mass of the atom?

Answer 2

In the nucleus

Question 3

Particle Relative Mass Relative Charge
Proton
Neutron
Electron

Answer 3

1 +1
1 0
negligible (5x10^-4) -1

Question 4

Definition of atomic number (Z)

Answer 4

the number of protons in the nucleus of an atom

Question 5

Definition of mass number (A)

Answer 5

the number of protons and neutrons in the nucleus of an atom.

Question 5

How are positive and negative ions formed? What are they called?

Answer 5

+ve loses electron(s) cation
-ve gains electron(s) anion

Question 5

Calculate % composition of Ir-191 and Ir-193. Atomic mass = 192.22.

Answer 5

[191x + 193(100 - x)] / 100 = 192.22
x = 39
Ir-191 = 39%
Ir-193 = 61%

Question 5

What equipment is used to measure relative abundance of isotopes?

Answer 5

Mass spectrometer

Question 5

Definition of isotope

Answer 5

Different atoms of the same element with different mass numbers (numbers of neutrons in the nucleus). Have the same chemical properties (same number of electrons) but different physical properties (melting and boiling points) (different masses so move at different speeds)

Question 5

Calculate Ar of Li-6 7% and Li-7 93%

Answer 5

[(6x7) + (7x93)] / 100 = 6.93

Question 6

What is the maximum number of electrons in the first five main energy levels?

Answer 6

1: 2
2: 8
3: 18
4: 32
5: 50

(2n^2)

Question 7

state the electromagnetic spectrum and increasing frequency, wavelength, energy.

Answer 7

radio wave, microwave, infrared, visible, UV, x-rays, gamma rays

–> increasing frequency
–> increasing energy
<– increasing wavelength

Question 8

What does a line spectrum for Hydrogen look like?

Answer 8

lines get closer at higher frequency. Further away from nucleus, higher energy levels.

Question 9

What is the difference between a line spectrum and a continuous spectrum?

Answer 9

Continuous spectrum: all energies / wavelengths
Line Spectrum: certain energies / wavelengths

Question 10

How does a spectrum arise? What does this give evidence of?

Answer 10

Electron promoted to higher energy level.
Becomes unstable and returns to lower energy level.
Gives out a photon of light, which gives a line in the spectrum.

Electrons being in energy levels – only allowed to have certain amounts of energy.

Question 11

What level does the electrons fall back to in: UV, visible, Infrared

Answer 11

UV: 1
visible: 2
Infrared: 3, 4, 5

Question 12

What region of the Spectrum does the electron fall back to energy level: 1, 2, 3, 4, 5

Answer 12

1: UV
2: Visible
3: Infrared
4: Infrared
5: Infrared

Question 13

What is the convergence limit?

Answer 13

When the lines merge to form a continuum. The electrons can have any energy and are free from influence of the nucleus (electrons no longer in atom)

Question 14

Give full electron configuration of As

Answer 14

1s^2, 2s^2, 2p^6, 3s^2, 3p^6, 4^2, 3d^10, 4p^3

Question 15

Give condensed electron configuration of Kr.

Answer 15

[Ar] 4s^2, 3d^10, 4p^6

Question 16

What is an orbital?

Answer 16

a region of space where there is a high probability of finding an electron (a discrete energy level)

Question 17

How many electrons can an orbital contain?

Answer 17

2

Question 18

What does the first shell (in orbitals) contain? What shape is it? Draw it.

Answer 18

1s orbital
spherical

Question 19

What does the second shell (in orbitals) contain? What shape are they? Draw them.

Answer 19

one 2s orbital, three 2p orbitals
dumb-bell shape – along the lines x, y, z.
degenerate (all the same energy)

Question 20

What does the third shell (in orbitals) contain? Draw them.

Answer 20

one 3s, three 3p, five 3d orbitals.

Question 21

What are the number of orbitals in each energy level?

Answer 21

s p d f
1: 1
2: 1 3
3: 1 3 5
4: 1 3 5 7
5: 1 3 5 7

Question 22

What is Pauli Exclusion Principle for electrons filling orbitals?

Answer 22

The maximum number of electrons in an orbital is two. If there are two electrons in an orbitals they must have opposite spins.

Question 23

What is Hund’s Rule for electrons filling orbitals?

Answer 23

Electrons fill orbitals of the same energy (degenerate orbitals) so as to give the maximum number of electrons with parallel spins.

Question 24

What are the electron configurations of 24Cr and 29Cu?

Answer 24

[Ar] 3d^5 4s^1
[Ar] 3d^10 4s^1

Question 25

What is the equation for the first ionisation energy for an element?

Answer 25

M (g) –> M^+ (g) + e^-

Question 26

What is the equation for the second ionisation energy for an element?

Answer 26

Mg^+ (g) –> Mg^2+ (g) + e^-

Question 27

Why is the second ionisation energy always higher than the first?

Answer 27

For the 1st IE, and e^- is removed from a neutral atom, but for the 2nd it is removed from a +ve ion. Takes more energy to remove a -vely charged e^- from a +ve ion.

Question 28

Why is the outer electron of potassium easy to remove?

Answer 28

it is shielded from the full attractive force of the nucleus by the other shells of electrons.

Question 29

Why is the next electron in potassium considerably more difficult to remove?

Answer 29

the electron is in a new shell and is closer to the nucleus and more strongly attracted, and only shielded by ten electrons (electrons in same shell don’t shield each other well) so more strongly attracted.

Question 30

What is the general trend of ionisation energy in elements? Why?

Answer 30

increases from left to right across a period.

Electrons added to same shell, but electrons in same shell don’t shield each other well. Force on outer electrons increases from left to right and outer electron is more difficult to remove for neon.

Question 30

Why does Boron have a lower first ionisation energy than beryllium?
Be: 1s^2 2s^2
B: 1s^2 2s^2 2p^1

Answer 30

2p sublevel in B is higher in energy than 2s sub level in Be. More energy required to remove e^- from Be.

p e^- in B is shielded to a certain extent by the s electrons so easier to move.

Question 31

Why is the first ionisation energy of oxygen lower than nitrogen?

2p 2p
| | | || | |
N O

Answer 31

No p e^- paired in N, less repulsion. 2 e^- paired in same p orbital in O. Electrons repel each other so easier to remove.

Question 32

How does ionisation energy change in a group? Why?

Answer 32

decreases down a group.

Smaller atoms higher up. Outer e^- is closer to the nucleus and more strongly attracted, so more energy needed to remove it.